Chemical Quantities

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Chapter 12

• Chemical Quantities

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How do you measure things?

• We measure mass in grams.
• We measure volume in liters.
• We count atoms or compounds in MOLES.

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Moles

• Defined as the number of carbon atoms in exactly
  12 grams of carbon-12.
• 1 mole is 6.02 x 1023 particles.
• 6.02 x 1023 is called Avogadros number.
• MEMORIZE this number!

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Types of questions

• How many molecules of CO2 are in 4.56 moles of
  CO2 ?
• How many moles of water is 5.87 x 1022
  molecules?
• How many atoms of carbon are there in 1.23 moles
  of C6H12O6 ?
• How many moles is 7.78 x 1024 formula units of
  MgCl2?

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Measuring Moles

• The decimal number on the periodic table is also
  the mass of 1 mole of those atoms in grams.
• Called molar mass
• on PT 1 mole

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Examples

• How much would 2.34 moles of carbon weigh?
• How many moles of magnesium in 24.31 g of Mg?
• How many atoms of lithium in 1.00 g of Li?
• How much would 3.45 x 1022 atoms of U weigh?

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What About Compounds?

• in 1 mole of H2O molecules there are two moles of
  H atoms and 1 mole of O atoms
• To find the mass of one mole of a compound
• determine the moles of the elements they have
• Find out how much they would weigh
• add them up

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What About Compounds?

• What is the mass of one mole of CH4?
• 1 mole of C 12.01 g
• 4 mole of H x 1.01 g 4.04g
• 1 mole CH4 12.01 4.04 16.05g
• The molar mass of CH4 is 16.05g

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Molar Mass

• The mass of one mole of a compound.
• What is the molar mass of Fe2O3?
• 2 moles of Fe x 55.85 g 111.70 g
• 3 moles of O x 16.00 g 48.00 g
• The molar mass 111.70 g 48.00 g 159.70g

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Examples

• Calculate the molar mass of the following
• Na2S
• N2O4
• C60
• Ca(NO3)2
• C6H12O6
• (NH4)3PO4

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Using Molar Mass

• Finding moles of compounds

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Molar Mass

• The number of grams of 1 mole of atoms, ions, or
  molecules.
• Make conversion factors
• Change grams of a compound to moles of a compound.

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For example

• How many moles is 5.69 g of NaOH?

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For example

• How many moles is 5.69 g of NaOH?

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

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For example

• How many moles is 5.69 g of NaOH?

• need to change grams to moles
• Need molar mass for NaOH
• 1mole Na 22.99g 1 mol O 16.00 g 1 mole of
  H 1.01 g
• 1 mole NaOH 40.00 g

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Examples

1. How many moles is 4.56 g of CO2 ?
2. How many grams is 9.87 moles of H2O?
3. How many molecules in 6.8 g of CH4?
4. 49 molecules of C6H12O6 weighs how much?

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Gases and the Mole
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Gases

• Many of the chemicals we deal with are gases.
• Difficult to weigh
• How do we know how many moles of gas we have?

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Gases

• Two things effect the volume of a gas
• Temperature
• Pressure
• Compare at the same temperature and pressure.

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Standard Temperature and Pressure

• STP is 0ºC and 1 atm pressure
• At STP 1 mole of gas occupies 22.4 L
• This is called molar volume

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Avogadros Hypothesis

• at the same temperature and pressure, equal
  volumes of gas have the same number of particles.

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Examples

• What is the volume of 4.59 mole of CO2 gas at
  STP?
• How many moles is 5.67 L of O2 at STP?
• What is the volume of 8.8g of CH4 gas at STP?

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Density of a Gas

• The units will be g / L
• We can determine the density of any gas at STP if
  we know its formula.
• If you assume you have 1 mole, then the mass is
  the molar mass (P.T.)
• At STP the volume is 22.4 L.

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Examples

1. Find the density of CO2 at STP.
2. Find the density of CH4 at STP.

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The other way

• Given the density, we can find the molar mass of
  the gas.
• Again, pretend you have a mole at STP, so V
  22.4 L.
• m is the mass of 1 mole, since you have 22.4 L of
  a gas.
• What is the molar mass of a gas with a density of
  1.964 g/L?
• 2.86 g/L?

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All the things we can change
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We have learned how to change

• moles to grams
• moles to atoms
• moles to compounds
• moles to liters
• compounds to atoms
• compounds to ions

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Mass
Moles
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Mass
PT
Moles
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Mass
Volume
PT
Moles
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Mass
Volume
22.4 L
PT
Moles
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Mass
Volume
22.4 L
PT
Moles
Compounds
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Mass
Volume
22.4 L
PT
Moles
6.02 x 1023
Compounds
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Mass
Volume
22.4 L
PT
Moles
6.02 x 1023
Compounds
Atoms
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Mass
Volume
22.4 L
PT
Moles
6.02 x 1023
Compounds
Ions
Atoms
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Percent Composition

• Part x 100
• Whole
• Find the mass of each component, then divide by
  the total mass.

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Example

• Calculate the percent composition of a compound
  that is 29.0 g of Ag with 4.30 g of S.

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Getting percent composition from the formula

• If we know the formula, assume you have 1 mole.

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Examples

• Calculate the percent composition of C2H4
• Calculate the percent composition of Aluminum
  carbonate.

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Empirical Formula

• From percentage to formula

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• Empirical Formula - The lowest whole number ratio
  of elements in a compound.
• Molecular Formula - the actual ratio of elements
  in a compound.

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Empirical and Molecular Formula

• - The two can be the same.
• CH2 empirical formula
• C2H4 molecular formula
• C3H6 molecular formula
• H2O both

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Calculating Empirical

• Find the lowest whole number ratio
• C6H12O6
• CH4O
• It is not just the ratio of atoms, it is also the
  ratio of moles of atoms.
• In 1 mole of CO2 there is 1 mole of carbon and 2
  moles of oxygen.
• In one molecule of CO2 there is 1 atom of C and 2
  atoms of O.

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Calculating Empirical

• We can get ratios from percent composition by
  assuming you have 100g.
• The percentages become grams.
• Can turn grams to moles.
• Find lowest whole number ratio by dividing by the
  smallest.

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Example

• Calculate the empirical formula of a compound
  composed of 38.67 C, 16.22 H, and 45.11 N.
• Assume 100 g so
• 38.67 g C x 1mol C 3.220 mole C
  12.01 gC
• 16.22 g H x 1mol H 16.09 mole H
  1.01 gH
• 45.11 g N x 1mol N 3.219 mole N
  14.01 gN

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Example

• The ratio is
• 3.220 mol C 1 mol C 3.219 molN
  1 mol N
• The ratio is
• 16.09 mol H 5 mol H 3.219 molN
  1 mol N
• C1H5N1

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Example

• Caffeine is 49.48 C, 5.15 H, 28.87 N and
  16.49 O. What is its empirical formula?

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Empirical to molecular

• empirical formula lowest ratio
• Actual molecule would weigh more by a whole
  number multiple.
• Divide the actual molar mass by the mass of one
  mole of the empirical formula.
• Caffeine has a molar mass of 194 g. what is
  its molecular formula?